Equilibrium
Physical Equilibrium An example of physical Equilibrium is when the number of water particles evaporating and the number of condensing particles become equal . Physical Equilibrium is the state in which two reverse process occur at the same rate . Chemical Equilibrium Chemical Equilibrium is a process in which two reverse reactions occur at the same rate . Thus , chemical equilibrium is possible only in reversible reactions . The reaction taking place from Reactants to Products is called Forward Reaction . The reaction taking place from Products to Reactants is called Reverse reaction . At equilibrium , the rate of Forward Reaction and Reverse reaction becomes same and the concentrations attain constant values . Thus , equilibrium is dynamic in nature , although the concentrations become static . Law of Mass Action The Law of Mass Action states that the rate of a Reaction is directly proportional to the products of their concentrations raised to the power of their respective stoichiometric coefficients . Equilibrium Constant Equilibrium Constant = Rate Constant of Forward Reaction / Rate Constant of Reverse Reaction Kp = KcdnRT Le Chatelier's Principle The Le Chatelier's Principle states that when a stress (additional change) is applied to a reacting system in equilibrium , the system shifts itself in such a way so as to reduce the strain caused by the stress . Significance of ΔG & ΔG0 = Ionic Equilibrium = Ionization is a reversible process .Thus Ionic equilibrium is the equilibrium associated with ionization . Electrolytes Electrolytes are those solutions through which electricity can be passed . Strong Electrolytes dissociate completely while weak electrolytes dissociate incompletely . Thus , the dissociation of weak electrolytes id=s a reversible process . Dissociation Degree of Dissociation α = No. of moles dissociated / Total No. of moles For strong electrolytes , the α = 1 . Acids and Bases Arrenhius Theory Acids are H+ donating species . Bases are OH- donating species Lowry - Brownsted Theory Acids are proton donating species . Bases are proton accepting species . Lewis Theory Acids are electron accepting species . Bases are electron accepting species . Strength of Acid and Base :-''' Ka = H+A- / HA = Cα 2 / 1 - α (Ostwald's Dilution Law) Kb = OH-B+ / BOH = Cα2 / 1 - α Cα2 is constant . Ka & Kb are just equilibrium constants and hence depend only on temperature . The strengths of acids and bases also depends upon the solvent used . If the solvent can accept protons easily , the acid will donate more protons . Concept of pH pH stands for Power of Hydrogen . It is the negative Logarithm of Concentration of H+ ions to the base 10 . pH = - log10H+ OR pKa = - log10Ka Since it is negative , lower the pH , greater is the power of hydrogen ; stronger is the acid pOH is the negative Logarithm of Concentration of OH- ions to the base 10 . pOH = -log10OH- OR pKb = - log10Kb pH + pOH = 14 '''Relative Acid Strength 1) Polarity of bond to which Hydrogen is attached to : 'The more polar this bond is , the more easily protons are given away . 2) Bond Strength ': '''The less the bond strength(bond with hydrogen) , the stronger the acid . i) HF < HCl < HBr < HI ii) H2O < HF iii) HIO < HBrO < HClO iv) HClO < HClO2 < HClO3 < HClO4 v) HSO4- < H2SO4 Ionic Product of Water For dissociation of Water , K = H+OH- / H2O K H2O= H+OH- Kw = H+OH- Hydrolysis of Salts Buffer Solutions Buffer solutions are the solutions whose pH does not change considerably with addition of small amounts of acid or base . Buffer solutions are generally prepared by mixing a weak acid(e.g. CH3COOH and CH3COONa) and it's salt or weak base and it's salt (e.g. NH4OH and NH4Cl). i.e. weak acid (CH3COOH) and it's conjugate base(CH3COO-) or weak base (NH4OH) and its conjugate acid(NH4+) . '''Henderson - Hasselbalch Equation Acidic Buffer : ''' Ka = H+A- / HA H+ = Ka HA / A- = Ka acid / salt '''pH = pKa - log10{acid / salt} A- is the concentration of salt , since it is completely dissociated . Basic Buffer : Kb = OH-B+ / BOH OH- = Kb BOH / B+ = Kb base / salt pOH = pKb - log10{base / salt} Conjugate Acids and Base A strong acid gives a weak conjugate base and a strong base gives a weak conjugate acid . Solubility Product For dissociation of AgCl , K = Ag+Cl- / AgCl AgCl = K' = constant KK' = Ag+Cl- Ksp = Ag+Cl- 1. If I.P. > S.P. ; Precipitate is formed 2. If I.P. = S.P. ; Equilibrium 3. If I.P. < S.P. ; No Precipitate formed Order of Solubility depends upon concentration of one of the Ions . Common Ion Effect Common Ion Effect states that when a weak electrolyte is dissolved in strong electrolyte having an ion common to each other ; the dissociation of the weaker electrolyte is suppressed by the strong electrolyte . Tips and Tricks # While dealing with Ksp , Ionic Product , ka and kb , Rely on the equations and it's stoichiometry . Category:Chemistry